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TOPIC 1 : INTRODUCTION TO CHEMISTRY

MEANING OF CHEMISTRY

Chemistry is the branch of science dealing with elements and the compounds they form and the

reactions they undergo.

BRANCHES OF CHEMISTRY

The braches of chemistry include physical, environmental, analytical, industrial, organic and

inorganic chemistry.

a. PHYSICAL CHEMISTRY

It is the study of how chemical compounds and their constituents react with each other.

b. ENVIRONMENTAL CHEMISTRY

It is the study of how chemicals react naturally in the environment and human impact on natural

systems.

c. ANALYTICAL CHEMISTRY

It is the study of separation, identification, and quantification of the chemical components of

natural and artificial materials.

d. INDUSTRIAL CHEMISTRY

It is the study of the application of physical and chemical processes towards the change of raw

materials into beneficial products.

e. ORGANIC CHEMISTRY

It is the study of compounds that contain carbon except oxides of carbon and carbonates.

f. INORGANIC CHEMISTRY

It is the study of compounds that do not contain carbon and non-living things.

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IMPORTANCE OF CHEMISTRY IN EVERYDAY LIFE

Chemistry is important in everyday life and it is applied in different ways. Some of the

applications of chemistry are:

 Water treatment. Different chemical processes are used to purify water so that it is safe for

drinking.

 Cooking nsima. Mixing of the ingredients applies concepts in chemistry.

 Making a cup of tea.

 Pharmaceuticals.

 Food industries. Chemistry is involved in the processing of the food. For example, lime is

added to brown sugar so that it becomes white.

 Manufacture of soap and detergents also applies knowledge of chemistry.

 Manufacture of pesticides.

AREAS WHERE CHEMISTRY IS APPLIED

 Pharmaceutical companies that manufacture medical drugs.

 Companies that make food and drinks (soft and alcoholic).

 Companies that manufacture oil products.

 Companies that manufacture fertilizers and pesticides.

 Water purification and supply companies.

 The mining industries.

CAREERS IN CHEMISTRY AND THEIR IMPORTANCE

Most careers in modern society require the application of the knowledge in chemistry.

a. Medicine and nursing

Doctors and nurses need chemistry as part of their training.

b. Pharmacist

Pharmacists require chemistry as part of their training in order to understand the chemicals they

are providing.

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c. Food chemist

Food chemists help test manufactured food to ensure that it is safe to eat.

d. Teacher chemist

Chemistry teachers prepare students for different careers by teaching them chemistry in schools.

e. Environmental chemist

Environmental chemists are involved in managing the environment by making sure rivers and

lakes do not becomes polluted, waste materials are properly disposed among others.

LABORATORY

A laboratory is a special room equipped for conducting scientific research and experimentation.

SAFETY RULES IN THE LABORATORY

 Do not drink, taste nor eat anything in the laboratory. Any chemical is never tasted in the

laboratory but can be tested.

 Handle all materials in the laboratory carefully. Glassware must be held with both hands.

 Never run or play in the laboratory.

 Wear protective materials such as lab coat, an apron, and safety goggles.

 Never work in the laboratory barefooted.

 Avoid disturbing or pushing a colleague who is busy working in the laboratory.

 Clean all equipment and workplaces after each laboratory period.

 Tie up long hair and avoid wearing loose clothing which could be caught in equipment.

 Follow experimental procedures and do not take short cuts.

 Avoid carrying out any other experiments other than the one given by the teacher.

 Turn off water, gas and electricity outlets when not in use.

 Keep flame and flammable solutions apart.

 Keep electrical equipment away from water and keep areas around electrical equipment dry.

 Always clean glassware before using them.

 Always work in a well ventilated area.

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 Keep hands away from face, eyes, mouth and body while using chemicals or laboratory

equipment.

 Work areas should be kept clean.

 When first entering a laboratory, do not touch any equipment or other materials in the

laboratory until you are instructed to do so.

COMMON LABORATORY APPARATUS

PROTECTIVE EQUIPMENT IN THE LABORATORY

In the chemistry, some of the protective equipment includes goggles, gloves, lab coats,

respiratory/gas mask, eye wash station and fire extinguisher

HAZARD SYMBOLS

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A hazard symbol tells you the dangers associated with handling laboratory chemical and

apparatus. Some of them are shown below.

Hazard symbol Meaning

Harmful or irritant substance

Toxic substance

Highly flammable

Corrosive

Dangerous to the environment

THE SI UNIT SYSTEM OF MEASUREMENT

The system of measurement used nowadays is known as the SI system of units. SI stands for

International System.

BASIC UNITS

Basic units are a set of unrelated units that form the basis of the SI system of units. These

quantities cannot be expressed in terms of other quantities.

Quantity Unit name Unit Symbol

Length Metre M

Mass Kilogram Kg

Time Second S

Temperature Kelvin K

DERIVED UNITS

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These are units obtained from combination of basic units through multiplication and division.

Some of them are shown below.

Quantity Name of unit Unit symbol

Area square metre m

Volume cubic metre m

Speed metre per second m/s

Density kilogram per cubic metre kg/m

SI PREFIXES FOR UNITS OF MEASUREMENTS

The SI prefix is an affix that is added to the name of a basic unit. It tells you whether the unit is a

multiple or a fraction of the basic unit. Examples of SI prefixes are shown in the table.

Prefix Fraction or multiple

nano



micro



mili



kilo

mega

giga

MEASURING PHYSICAL QUANTITIES

Physical quantities must be measured as accurately as possible. Different measuring instruments

are used to measure the physical quantities.

a. MEASURING MASS

The mass of an object is measured using a balance. The triple beam balance is the commonly

used balance in the chemistry laboratory. It is shown in the figure below.

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HOW TO USE THE TRIPPLE BEAM BALANCE

 Put the balance on a flat hard surface where no wind is blowing.

 Put all the masses to zero marks.

 Adjust the zeroing screw until the pointer is at zero.

 Place the object whose mass you want to measure on the pan carefully.

 Move the big mass (100g) first, gradually until the beam balance topples over. Then push the

mass one step back.

 Move the second mass (10g) gradually until the beam balance topples over again and push it

one step back.

 Move the smaller mass (1g) gradually until the pointer is at zero mark again.

 Read the mass by adding the numerical values of positions of all the three masses. An

example is shown.

The mass of the object = (300 + 40 + 5) = 345 g

b. MEASURING VOLUME

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The instrument for measuring volume of a liquid is the measuring cylinder. The measuring

cylinders are in different sizes such as 25 ml, 50 ml, 100 ml, 500 ml, and 1000 ml.

HOW TO READ THE VOLUME

The volume of the liquid in the measuring cylinder must be taken from the bottom of meniscus.

The meniscus is the curve on the liquid caused by the liquid being slightly attracted to the glass.

The meniscus must be viewed at eye level and not from an angle as shown below.

c. MEASURING TEMPERATURE

Temperature is measured using an instrument called thermometer. There are different kinds of

thermometers but the commonly used in the chemistry laboratory is the liquid–in–glass

thermometer.

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This kind of thermometer has a glass tube sealed at both ends and is partly filled with liquid like

mercury or alcohol. When the liquid in the bulb is heated, it expands and the top of the liquid

moves up the tube.

To measure the temperature of a liquid, suspend the tip of the thermometer below the surface of

the liquid. Allow the liquid in the thermometer to expand. When it has finished expanding and it

is no longer moving up the column, you can read the thermometer.

d. MEASURING TIME

Time is measured using an instrument called stopwatch. A stopwatch is handheld and come in

different versions. The most commonly used in the chemistry laboratory is the digital stop watch.

Timing functions are controlled by two buttons; start/stop button and the reset button. Pressing

the start/stop button starts the timer running, and pressing it a second time stops it. Pressing the

reset button resets the stopwatch to zero.

When you begin a task, you press the start/stop button to begin recording your time. When you

finish working on that task, press the button again to stop recording time.

When the stop watch is set to zero, five zeroes will be displayed on the screen. The first two

zeroes on the left represent the “minutes”, the middle zeroes represent “seconds” and the last

zero “hundredths of a second”.

To read the time taken for the task, combine the numbers for your full time used. For example,

11:14:01 would be 11 minutes, 14 seconds and 01 hundredths of a second.

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SCIENTIFIC METHOD OF INVESTIGATION

Like other subjects, chemistry is a practical subject. It follows the scientific method or enquiry.

STEPS OF SCIENTIFIC INQUIRY

1. Identify a problem

Identifying a problem involves asking questions about the natural world. Examples of scientific

questions are:

 What causes rusting?

 Why do plastics not decompose easily?

2. Formulating a hypothesis

A hypothesis is a guessed answer to a problem. A hypothesis is formulated from the scientist’s

experiences and knowledge.

3. Testing the hypothesis

To test, the hypothesis, an experiment is carried out. An experiment is a series of investigations

intended to accept, modify or reject a hypothesis.

4. Analyse the results

This involves looking at the collected date and making sense of it.

5. Conclusion

The conclusion is drawn based upon the collected data. It is either a conformation or the

rejection of the hypothesis under investigation. If the hypothesis is correct, it is confirmed and

adopted and if false it is declared null and void hence rejected. When the hypothesis is rejected

another one is formulated and tested.

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TOPIC 2 : ESSENTIAL MATHEMATICAL SKILLS IN CHEMISTRY

EXPRESSING NUMBERS IN STANDARD FORM

The standard form or the scientific notation is a special way of writing very large number or very

small numbers. When a number is expressed in standard form, its meaning or value does not

change.

The number is written in two parts which give the original number when multiplied. One of the

numbers must be between 1 and 10 and the other is a power of ten.

EXPRESSING VERY BIG NUMBERS IN STANDARD NOTATION

Place a decimal point just after the first digit, followed by



10 to the power the number of places

moved from the decimal point to the last digit. The movement is from the left to the right. The

power of ten must be positive.

Example

Write down each of the following numbers in standard form.

a. 4 500

b. 67 413

c. 300 000 000

Solution

4.5×10

6.7413×10

3.0×10

EXPRESSING VERY SMALL NUMBERS IN STANDARD FORM

Place the decimal point just after the first non-zero digit, followed by



10 to the power the

number of places moved from the decimal point. The movement is from the right to the left. The

power of ten is negative.

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Example

Write the following numbers in scientific form.

a. 0.00067

b. 0.00145

c. 0.335

Solution

6.7 10





1.45 10





3.35 10





SIGNIFICANT FIGURES

Significant figures of a number are those digits that carry meaning contributing to its precision.

GUIDELINES FOR WRTING SIGNIFICANT FIGURES

1. All non-zero digits are significant. For example in the number 6753 there are 4 significant

figures.

2. Zeroes between non-zero digits are significant. For example, the number 40072 has 5

significant figures.

3. Zeroes to the left of non-zero digits are not significant. For example the number:

0.89 has 2 significant figures

0.089 has 2 significant figures

0.0089 has 2 significant figures

4. If a number ends in zeroes to the right of a decimal point, the zeroes are significant. For

example the number 9.0 has 2 significant figures.

5. In a figure without a decimal point, the right most non-zero digit is the least significant

figure. For example in 7900 the least significant is 9.

6. If the next digit after the last significant figure is 4 or less, the number is rounded down. If it

is 5 or more, it is rounded up. For example 14.628 to 4 significant figures is 14.63 while

15.473 to 4 significant figures is 15. 47.

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EXPRESSING NUMERICAL RESULTS TO CORRECT NUMBER OF SIGNIFICANT

FIGURES

a. Addition and subtraction

The result obtained from adding or subtracting and multiplying or dividing numbers must be

quoted based on the number that has the least number of significant figures.

Example 1

Add the following number: 2 345, 7 800 and 934 456.

Solution

The answer is 940 000 and not 944 601 because 7 800 has the least number of significant figures

(which is 2).

Example 2

Work out

2.467 465

2.7



Solution

424.872222222 to 2 significant figures is 420.

ACCURACY AND PRECISION

Accuracy is how close a measured value is to the actual value. Precision is how close the

measured values are to each other.

Example

Consider the three measurements 30.01g, 30.02 and 30.03. If the actual value is 30.0g, then all

the three measurements are accurate because they are very close to the actual value. They are

also precise because they are close to each other.

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GRAPHS

Graphs are pictorial representations of data values measured in an experiment.

TYPES OF GRAPHS

Graphs can be in form of line graphs, bar graphs or pie charts.

a. LINE GRAPHS

Line graphs uses points connected by a line to show data. The line graph must have the following

components:

 Title. For example “A graph of temperature against time”

 Axes. These are fixed reference lines for the measurement of coordinates. There are two axes

on the graph. The horizontal axis, a straight line drawn from the left to the right and the

vertical axis, a straight line drawn from bottom to the top. From any statement, “a graph of A

against B”, A must be on the vertical axis while B must be on the horizontal axis.

 Scale. A range of numbers that show the units used on the graph.

 Origin. It is a point where the vertical and horizontal axes meet.

An example of a line graph is shown below

b. BAR GRAPHS

A bar graph displays data using bars or rectangles to show comparisons between categories of

data. The bars can either be vertical or horizontal. On the bar graph, one axis will describe the

types of categories being compared; and another will have numerical values that represent the

values of data.

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Just like the line graph, the bar graph includes a title, scale, axes among other things. For

example, the figure below is a bar graph showing concentration of solute of three

chromatograms, A, B and C.

c. PIE CHARTS

A pie chart display data in the form of a circle. In the pie chart, a circle is divided into various

sections or segments. Each segment represents a certain proportion or percentage of the total. In

such a diagram, the total of all the given items is equated to 360º. The degrees of angles,

representing different items are calculated proportionately. An example of a pie chart is shown

below.

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TOPIC 3: COMPOSITION AND CLASSIFICATION

OF MATTER

MATTER

Matter is defined as anything that has mass and occupies space.

STATES OF MATTER

There are three states of matter. These are solids, liquids and gases

PROPERTIES OF MATTER

1. SOLIDS

 Particles in solids are tightly packed, usually in a regular pattern.

 They do not flow. This is because particles are not free to move. But they vibrate in fixed

positions.

 They have a definite shape.

 They have a definite volume.

 They are difficult to compress.

2. LIQUIDS

 Particles in liquids are close together, but with no regular arrangement.

 Liquids flow. This is because their particles are free to move while sticking together.

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 They have indefinite shape. In other words, they take the shape of the container in which

they are put.

 They have a definite volume.

 They are difficult to compress.

3. GASES

 Particles are very far apart from each other with no regular arrangement.

 Particles move randomly at very high speeds in all directions.

 They have an indefinite shape.

 They an indefinite volume.

 They can be compressed.

 They flow and particles can move past one another.

THE PARTICULATE NATURE OF MATTER

The particulate nature of matter is an idea that explains how matter is put together.

Matter is made up small particles called molecules. The molecules are in turn made up of

indivisible and invisible particles called atoms.

ATOM

An atom is defined as the smallest particle of matter.

EVIDENCE OF THE PARTICULATE NATURE OF MATTER

The particulate nature of matter can be proved using the concept of diffusion.

DIFFUSION

Diffusion is defined as the movement of particles from a region of a higher concentration to a

region of a lower concentration.

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Diffusion takes place mainly in liquids and gases, but it occurs most quickly in gases because

the particles in gases are very far apart and move more randomly at very high speeds.

It does not take place in solids because particles in solids are not free to move.

INVESTIGATING DIFFUSIONS IN LIQUIDS

Materials:

 A beaker

 Water

 Potassium permanganate crystals

 Thistle funnel.

Procedure

a. Put the thistle funnel into the beaker.

b. Slide a few crystals of potassium permanganate into the beaker through the thistle funnel.

c. Pour water into the beaker with the thistle funnel in the same position as shown in the

Figure 3.1.

Figure 3.1

d. Carefully, remove the thistle funnel.

e. Leave the set up for 5 minutes and observe what happens. Do not shake or swirl the

beaker.

Observations

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After five minutes, the purple colour of potassium permanganate is distributed throughout the

water in the beaker.

Explanations

The purple colour of potassium permanganate spreads because the particles leave the crystals

and mix with water. The solute particles move throughout the water as they keep colliding with

water particles as they move, eventually becoming evenly distributed and the colour of water

turns purple. In simple words, the particles of potassium permanganate have diffused.

This demonstration proves that matter is made up of tiny particles which are in constant motion.

COMMON EXAMPLES OF DIFFUSION IN EVERYDAY LIFE

 When you enter a restaurant, you smell food being cooked.

 Sugar diffuses in water when it dissolves.

 Coffee grains spread in water.

 A person wearing perfumed clothes

ELEMENTS

An element is a substance which cannot be split into simpler substances by chemical means.

Elements are made up of atoms. Some elements exist as separate atoms, while others comprise

groups of atoms combined together.

The groups of atoms that result from combination of two or more atoms are called molecules.

However, it is worth noting that elements are made form only one kind of atoms.

EXAMPLES OF ELEMENTS

There are about 115 elements which have been discovered. Ninety–nine of the elements occur

naturally, while twenty–four of these have been made artificially by scientists. The table shows is

a list of the first twenty elements.

Hydrogen Sodium

Helium Magnesium

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Lithium Aluminum

Beryllium Silicon

Boron Phosphorous

Carbon Sulphur

Nitrogen Chlorine

Oxygen Argon

Fluorine Potassium

Neon Calcium

CHEMICAL SYMBOLS OF ATOMS OF ELEMENTS

The chemical symbol is a shorthand form for writing the names of the elements. The system of

writing symbols uses the letter taken from the name of the element. This could be the English or

Latin name of the element.

RULES FOR WRITING CHEMICAL SYMBOLS OF ELEMENTS

 All chemical symbols consist of one or two letters.

 The second letter is added where some elements have the same initial letter.

 The first letter of the chemical symbols must always be capital.

Name of element Chemical symbol Name of element Chemical symbol

Hydrogen H Sodium Na

Helium He Magnesium Mg

Lithium Li Aluminium Al

Beryllium Be Silicon Si

Boron B Phosphorous P

Carbon C Sulphur S

Nitrogen N Chlorine

Oxygen O Argon Ar

Fluorine F Potassium K

Neon Ne Calcium Ca

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The table below shows chemical symbols of elements which are derived from Latin names of the

elements.

Element Latin name Chemical symbol

Sodium Natrium Na

Potassium Kalium K

Copper Cuprum Cu

Iron Ferrum Fe

Silver Argentum Ag

Lead Plumbum Pb

Gold Aurum Au

Mercury Hydrargyrum Hg

MOLECULES

A molecule is the smallest particle of an element or a compound which can exist in a free and

separate state.

TYPES OF MOLECULES

There are three main types of molecules. These are:

a. MONOATOMIC MOLECULES

These are molecules composed of one atom. Examples are helium (He), Neon (Ne), Argon (Ar),

Xenon (Xe) and Radon (Rn).

b. DIATOMIC MOLECULES

These are molecules composed of two atoms. Examples are Hydrogen (H

), Oxygen (O

Nitrogen (N

), Fluorine (F

), Chlorine (C

), Bromine (

) and Iodine (

c. POLYATOMIC MOLECULES

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These are molecules composed of many atoms. The e amples are Sulphur (

) and Phosphorous

(

USING MODELS OR DIAGRAMS TO ILLUSTRATE FORMATION OF MOLECULES

Example 1: Oxygen molecule (O

)

Let repressent one oxygen atom.

Since oxygen molecule is diatomic, it will be shown as:

oxygen molecule

Example 2: Water molecule (H

A water molecule consists of two hydrogen atoms and one oxygen atom.

Let represent oxygen atom

and represent hydrogen atom

The water molecule is shown as

water molecule

COMPOUNDS

A compound is a pure substance made up of two or more different chemical elements. In a

compound, elements are present in definite proportion. A compound has different chemical and

physical properties from those elements of which it is composed of.

The table below shows examples of compounds and elements which are present in each

compound.

Name of compound Elements present in the compound

Water Hydrogen, Oxygen

Carbon dioxide Carbon, oxygen

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Sodium chloride Sodium, Chlorine

Copper carbonate Copper, Carbon, Nitrogen

Glucose Carbon, Hydrogen , Oxygen

Ammonia Hydrogen, Nitrogen

Potassium nitrate Potassium, Nitrogen, Oxygen

Methane Carbon, Hydrogen

DIFFERENCES BETWEEN AN ELEMENT AND A COMPOUND

Element Compound

Is made up of only one kind of atoms Is made up of two or more elements

Cannot be split into simpler substances by

chemical means

Can be split into individual elements using

chemical reactions

Are represented by chemical symbols of their

atoms

Are represented by chemical formula

CHEMICAL FORMULAE OF SUSBTANCES

This is a formula that shows what types of atoms present and how many there are for each type,

in a given compound.

 The types of atoms present in the compound are indicated by the chemical symbols of the

elements.

 The numbers for each type of atoms is are indicated by a subscript. A subscript is a lowered

small digit. It is written slightly below the element symbol.

For example, the chemical formula of glucose is

6 1 2 6

C H O

. This formula tells us that:

 In glucose molecule, there are three types of atoms that combine. These are Carbon (C),

Hydrogen (H) and Oxygen (O).

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 There are 6 carbon (C) atoms, 12 Hydrogen (H) atoms and 6 Oxygen (O) atoms in the

compound.

The chemical formulae of some common substances are shown below.

Name of substance Chemical formulae

Carbon dioxide

Water

H O

Ethanol

2 5

C H OH

Sodium chloride (common salt)

NaC

Copper (II) sulphate

CuSO

DETERMINING THE TYPE AND NUMBER OF ATOMS IN THE CHEMICAL

FORMULA OF A GIVEN SUBSTANCE

Example 1:

The molecule of sucrose consists of 12 carbon atoms, 22 hydrogen atoms and 11 oxygen atoms.

Write down the molecular formula of sucrose.

Solution:

The molecular formula of sucrose is

12 22 11

C H O

Example 2:

The chemical formula of sodium sulphate is given as

2 4

Na SO

a. Determine the types and number of atoms in sodium sulphate.

b. Work out the total number of atoms in sodium sulphate.

Solution:

a. The types and number of atoms are:

Sodium (Na) = 2 atoms , Sulphur (S) = 1 atom, Oxygen (O) = 4 atoms

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b. Total number of atoms = 2 + 1 + 4 = 7 atoms

Example 3:

Work out the total number of atoms present in one molecule of copper nitrate,

3 2

Cu(NO )

Solution:

3 2

(NO )

means that everything inside the brackets is multiplied by 2.

Copper (Cu) = 1 atom

Nitrogen (N) = 2 atoms

Oxygen (O) = 6 atoms

Total number of atoms = 1 + 2 + 6

= 9 atoms

PURE SUBSTANCES AND MIXTURES

A pure substance is a substance that has definite and constant properties throughout the sample.

It can either be an element or a compound. Examples of pure substances are rain water, salt,

ethanol, sugar etc.

A mixture is a substance that contains two or more kinds of substances that are not chemically

joined together.

TYPES OF MIXTURES

There are two types of mixtures. These are homogeneous and heterogeneous mixtures.

a. HOMOGENEOUS MIXTURE

It is a mixture in which the particles are uniformly distributed throughout the mixture. Examples

are sugar solution, salt solution, blood, air, milk etc.

b. HETEROGENOUS MIXTURE

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It is a mixture in which the particles are not evenly distributed within the sample. Examples are

rock, soup, water, crude oil, a mixture of slat and sand.

SOLUTIONS

A solution is a homogeneous mixture of two or more components. It is a mixture of a solute and

solvent.

A solute is substance that dissolves in another substance. A substance in which another

substance dissolves is called the solvent. For example, in a sugar solution, sugar is the solute

while water is the solvent.

AQUEOUS SOLUTION AND NON-AQUEOUS SOLUTIONS

An aqueous solution is the one in which the solvent is water. Examples are sugar solution, salt

solution etc. A non-aqueous solution is the one in which the solvent is not water. Examples are

iodine in water, acetone in benzene.

TYPES OF SOLUTIONS

a. Solid–in–solid

This is a solution in which a solid is dissolved in another solid.

This is done when both metals are melted so that they can uniformly form a single solution. For

example, galvanized iron is a solution of zinc in iron, and brass is a solution of zinc in copper.

b. Liquid–in–liquid

This is a solution in which a liquid is dissolved in another liquid. For example, vinegar is a

solution of acetic acid in water.

c. Solid–in–liquid

This is a solution in which a solid dissolves in a liquid. The best example is a solution of sugar

in water.

d. Gas–in–liquid

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This is a solution in which a gas is dissolved in a liquid. Fizzy drinks are solutions of a gas in a

liquid. For example, in soda water, coca-cola, cocopina, and cherry plum carbon dioxide is

dissolved in water.

SATURATED AND UNSATURATED SOLUTIONS

A saturated solution is the one in which no more solute can dissolve. If more solute is added to

a saturated solution, undissolved crystals of the solute will rest at the bottom of the container

without dissolving. An unsaturated solution is the one in which more solute can dissolve.

WAYS OF MAKING A SATURATED SOLUTION UNSATURATED

 Adding more solvent to the solution.

 Increasing the temperature of the solution i.e. by heating

WAYS OF MAKING UNSATURATED SOLUTION SATURED

 Add more solute to a liquid.

 Evaporate the solvent from an unsaturated solution.

 Lowering the temperature of the solution.

FACTORS AFFECTING SOLUBILITY

a. TEMPERATURE

The solubility of solids in liquids increases as the temperature increases, and decreases as the

temperature decreases. When temperature increases, solvent particles gain kinetic energy and

collide frequently with solute particles, hence speeding up the solubility.

b. SIZE OF PARTCILES

Solubility of solids increases with an increase in surface area and decreases with a decrease in

surface area. Thus, powders dissolve more quickly than lumps in the same volume of solvent.

This is because; small particles can easily come into contact with solvent molecules, while it is

difficult for solvent molecules to surround large particles.

c. POLARITY

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Polar solutes dissolve best in polar solvents. Non–polar solutes dissolve best in non–polar

solvents.

WAYS OF INCREASING SOLUBILITY OF SOLUTES IN SOLVENTS

 Stirring the solution.

 Crushing the solute to powder.

 Heating the solution.

METHODS OF SEPARATING MIXTURES

The methods include filtration, decantation, evaporation, distillation (simple and fractional),

chromatography and magnetism.

a. FILTRATION

Filtration is a method used to separate an insoluble solid from a liquid. During filtration, the

mixture is passed through a filter paper which acts like a sieve. The filter paper has millions of

tiny holes in it, allowing the liquid to pass but retaining the solid particles.

The liquid that passes through the filter paper is called the filtrate and the solid that remains on

the filter paper is called the residue.

Examples of mixtures that can be separated by filtration include a mixture of sand water, and a

mixture of salt and a salt.

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b. DECANTATION

This method is used to separate mixtures of immiscible liquids or of a solid and a liquid in a

suspension.

Immiscible liquids are liquids that do not mix and form two distinct layers in a container e.g.

water and paraffin, vinegar and oil etc.

To separate two immiscible liquids, the top layer is carefully poured out into another container.

The middle portion is rejected since it is a mixture of the two unseparated liquids. The other

liquid remains behind.

When separating a mixture of a solid and a liquid, the mixture is allowed to stand in a beaker

until all the solid settles at the bottom of the container. Then, the liquid is carefully poured off to

leave the insoluble solid behind.

c. EVAPORATION

This is a method of separating a soluble solid substance from its solvent. For example, salt could

be recovered from salt solution using this method. Heat is applied to an evaporating basin

containing the solution and molecules of a liquid escape from it leaving behind the solid

particles.

Figure 2.3

d. DISTILLATION

This is a chemical process of separating substances that have different boiling points by heating

the mixture until it turns into gas and then cooling and collecting each substance separately.

Page | 30

There are two kinds of distillation; simple distillation and fractional distillation.

i. SIMPLE DISTILLATION

Simple distillation is used to separate a solvent from a solution. It is useful for producing water

from salt solutions.

Simple distillation works on the principle that the dissolved solute has a much higher boiling

point than the solvent. When the solution is heated, solvent vapour evaporates from the solution.

The gas moves away and is cooled and condensed. The remaining solution becomes more

concentrated in solute as the amount of solvent in it decreases.

ii. FRACTIONAL DISTILLATION

Fractional distillation is used to separate different liquids from a mixture of miscible liquids. It

is useful for separating ethanol from a mixture of ethanol and water, and for separating different

fractions from crude oil (petroleum).

Fractional distillation works on the principle that the different liquids have different boiling

points. When the mixture is heated:

 vapours rise through a column which is hot at the bottom and cold at the top. It is called the

fractionating column.

Page | 31

 vapours condense when they reach a part of the column that is below the temperature of

their boiling point.

The different liquids are collected from different parts of the column. The substance with the

lowest boiling point is collected at the top of the column.

SEPARATING ETHANOL FROM A MXITURE OF ETHANOL AND WATER

 The apparatus is set up as shown below.

 The mixture is heated and a mixture of ethanol vapour and water vapour rises up the column.

 The vapour condenses on the glass beads in the column making them hot.

 When the temperature of the beads reaches about 78ºC, ethanol vapour moves up the column

into the condenser, while the water drips back into the flask.

 Eventually, the thermometer reading rises above 78ºC. This indicates that all the ethanol has

been separated hence heating can be stopped.

Page | 32

e. CHROMATOGRAPHY

This is a method of separating a mixture of chemically similar ingredients. These are often

coloured substances such as food colourings, inks, dyes or plant pigments.

Chromatography separates mixtures by taking advantage of their different rates of movement in

a solvent over an absorbent material.

PAPER CHROMATOGRAPHY

In paper chromatography, the absorbent material is a filter paper and solvents such as water,

ethanol, propanone and other organic solvents can be used. Substances are picked up and

carried by a mobile phase which moves through a stationary phase.

The stationary phase is the part of the chromatography which is in solid state e.g. the filter

paper.

The mobile phase is the part of the chromatography which is in gaseous or liquid state e.g. the

solvent.

The different dissolved substances in the mixture are attracted to the two phases in different

proportions. This causes them to travel at different rates through the paper.

PROCEDURE

 Using a pencil, a ‘start line’ is drawn near the bottom of the chromatography paper.

 The mixture to be separated is spotted on the start line.

 The bottom of the chromatography paper is dipped into the solvent and the solvent travels up

the paper by capillary action.

 The solvent picks up the substance being separated and carries it up the paper.

 The different components in the substance rise to different heights.

 The “solvent front” is marked.

The solvent front is the furthest point reached by the solvent on chromatography paper.

The pattern formed by the substances that have been separated by chromatography is called the

chromatogram.

Page | 33

f. MAGNETSIM

This is a method used to separate magnetic from non-magnetic materials.

A magnet is hanged over the mixture and only magnetic materials are attracted to the magnet

thereby achieving separation.

Page | 34

TOPIC 4 : ATOMIC STRUCTURE

ATOM

An atom is the smallest particle of matter.

COMPOSITION OF AN ATOM

An atom consists of three sub-atomic particles. These include protons, neutrons and lectrons.

The central part of the atom is called the nucleus.

THE NUCLEUS

 It contains protons and neutrons.

 It has a positive charge because of the protons.

 The whole mass of the atom is concentrated in the nucleus.

 It is tiny compared to the size of the atom.

Outside the nucleus are regions called energy levels, or shells.

ENERGY LEVELS OR SHELLS

 These are imaginary paths through which electrons move.

 They are different distances from the nucleus.

 They are associated with a certain quantity of energy. An electron moving from one energy

level to another absorbs or emits energy depending on whether it moves to a lower or a

higher energy level.

Page | 35

CHARACTERISTICS OF PARTICLES THAT MAKE AN ATOM

1. PROTONS

 These are positively charged particles. Each proton has a charge of

1

 Each proton has a mass of 1 amu (atomic mass unit).

 They are found inside the nucleus of an atom.

2. ELECTRONS

 These are negatively charged particles. Each electron has a charge of

1

 Electrons have no mass.

 They move around the nucleus at very high speeds.

 They occupy the energy levels.

3. NEUTRONS

 These particles have no charge. They are considered to be neutral because they contain an

equal number of positive and negative charges.

 Each neutron has a mass of 1 amu.

 They are found inside the nucleus of an atom.

The table below shows the summary of characteristics of particles found in atom.

Particle Charge Mass Location

Proton +1 1 amu nucleus

Electron -1 Almost zero energy levels/shells

Neutron 0 (No charge) 1 amu nucleus

ELECTRON CONFIGURATION

Electron configuration is the number and arrangement of electrons in atom. In other words, the

electron configuration shows us how electrons are distributed in the energy levels of an atom.

The first twenty elements follow the following electron configuration:

 The 1

shell can hold a maximum of 2 electrons.

 The 2

shell can hold a maximum of 8 electrons.

 The 3

shell can hold a maximum of 8 electrons.

 The 4

shell can hold a maximum of 18 electrons.

The electron configuration is shown as a pattern of numbers. For example, 2.8.8.2 or 2:8:8:2.

WORKED EXAMPLES

ON ELECTRON CONFIGURATION

Example 1:

A sodium (Na) atom contains 11 atoms.

Write down its electron configuration.

Draw the atomic structure of sodium (Na)

Solution:

a. 1

shell =

2 electrons

From the 11 electrons, the remainder is 9.

shell =

8 electrons

From the 9 electrons, the remainder is 1.

shell = 1

Therefore, the electron configuration of sodium is 2.8.1

The atomic structure of sodium is shown as

shell can hold a maximum of 8 electrons.

shell can hold a maximum of 18 electrons.

The electron configuration is shown as a pattern of numbers. For example, 2.8.8.2 or 2:8:8:2.

ON ELECTRON CONFIGURATION

A sodium (Na) atom contains 11 atoms.

Write down its electron configuration.

Draw the atomic structure of sodium (Na)

2 electrons

From the 11 electrons, the remainder is 9.

8 electrons

From the 9 electrons, the remainder is 1.

Therefore, the electron configuration of sodium is 2.8.1

The atomic structure of sodium is shown as

Page | 36

The electron configuration is shown as a pattern of numbers. For example, 2.8.8.2 or 2:8:8:2.

Example 2:

Boron (B) has 5 electrons.

Write down its electron configuration.

Draw the atomic structure of boron.

Solution

Its electron configuration is 2.3

b. Its atomic struct

ure is shown as:

ATOMIC NUMBER AND MASS NUMBER

Atomic number

is the number of

letter Z

. The atomic number is like an ‘identity card’ of an atom. I

atoms of different elements different.

Mass number

is the number of protons and neutrons in the nucleus of

capital letter A.

The number of neutrons is normally represented by capital letter

The relationship between the atomic number (Z), the mass number (A) and the number of

neutrons (N) is expressed as follows:

Mass number = Numbe

r of protons +

A = Z

Write down its electron configuration.

Draw the atomic structure of boron.

Its electron configuration is 2.3

ure is shown as:

ATOMIC NUMBER AND MASS NUMBER

is the number of

protons

in the nucleus of an atom. It is denoted by capital

. The atomic number is like an ‘identity card’ of an atom. It is the property that makes

atoms of different elements different.

is the number of protons and neutrons in the nucleus of

an atom. It is denoted by

The number of neutrons is normally represented by capital letter

The relationship between the atomic number (Z), the mass number (A) and the number of

neutrons (N) is expressed as follows:

r of protons +

Number of neutrons

+ N

Page | 37

in the nucleus of an atom. It is denoted by capital

t is the property that makes

an atom. It is denoted by

The relationship between the atomic number (Z), the mass number (A) and the number of

Page | 38

The number of protons in an electrically neutral atom is equal to the number of electrons. The

atomic numbers, number of neutrons and mass numbers of the first twenty elements are shown

below.

Element Symbol Number of protons

(Atomic number, Z)

Number of

neutrons (N)

Mass number (A)

Hydrogen H 1 0 1

Helium He 2 2 4

Lithium Li 3 4 7

Beryllium Be 4 5 9

Boron B 5 6 11

Carbon C 6 6 12

Nitrogen N 7 7 14

Oxygen O 8 8 16

Fluorine F 9 10 19

Neon Ne 10 10 20

Sodium Na 11 12 23

Magnesium Mg 12 12 24

Aluminium Al 13 14 27

Silicon Si 14 14 28

Phosphorous P 15 16 31

Sulphur S 16 16 32

Chlorine

17 18 35

Argon Ar 18 22 40

Potassium K 19 20 39

Calcium Ca 20 20 40

NUCLIDE SYMBOLS

Page | 39

A nuclide is any form of an element which is characterized by specific constitution of its

nucleus.

To represent a nuclide, we write the chemical symbol of the element with the mass number on

the top left and atomic number at the bottom left of the element.

where A is the mass number, Z is the atomic number and X is the element symbol.

Example

You are given an atom X that has 17 protons, 20 neutrons and 17 electrons.

a. Find the atomic number of atom X.

b. Work out its atomic mass.

c. Write down the nuclide of X.

d. Write down the electron configuration of X.

Solution:

a. Atomic number = Number of protons = 17

b. Mass number = Number of protons + Number of neutrons

= 17 + 20

= 37 amu

c. The nuclide of X is

d. Since there are 17 electrons, the electron configuration of X is 2.8.7.

ISOTOPES

Isotopes are atoms of the same element that have the same atomic numbers but different mass

numbers.

Page | 40

They have different mass numbers because they have different number of neutrons on their

nuclei. Examples of isotopes are shown below.

Element Name of isotope Notation Particles present

Hydrogen Hydrogen – 1

1 proton, 0 neutrons

Hydrogen – 2

1 proton, 1 neutron

Hydrogen – 3

1 proton, 2 neutrons

Carbon Carbon – 12

6 protons, 6 neutrons

Carbon – 14

6 protons, 8 neutrons

Chlorine Chlorine – 35

17 protons, 18 neutrons

Chlorine – 37

17 protons, 20 neutrons

Uranium Uranium – 235

235

92 protons, 143 neutrons

Uranium – 238

238

92 protons, 146 protons

KEY POINTS ABOUT ISOTOPES

 They have the same number of protons on their nuclei.

 They have different number of neutrons on their nuclei.

 They have the same number of electrons in their outermost shell.

 They belong to the same group of the Periodic Table.

 They belong to the same period in the Periodic Table.

 They have identical (similar) chemical properties i.e. they react in the same way.

CALCULATING THE AVERAGE MASS OF AN ELEMENT GIVEN MASSES OF

ISOTOPES

The average mass of atoms is also called the relative atomic mass (RAM). It is calculated from

the relative percentage abundance in nature (RPA) of the isotopes and the relative isotopic mass

(RIM). Thus:

RAM of element X = (RPA of isotope 1



its RIM) + (RPA of isotope 2



its RIM)

Page | 41

Example 1:

A sample of chlorine gas contains two isotopes,

and

. Given that the percentage

abundance of

is 75% and and that of

is 25%, calculate the relative atomic mass of

chlorine.

Solution

Relative atomic mass of chlorine =

×35

100

×37

100

= 26. 25 + 9.25

= 35.5 amu

Example 2:

A sample of neon contains 90% of atoms of

N e

and 10% of atoms of

. Work out the

relative atomic mass of this sample of neon.

Solution

Relative atomic mass of neon =

× 20

100



= 18 + 2. 2

= 20. 2 amu

Page | 42

TOPIC 5 : THE PERIODIC TABLE

The periodic table is a table in which chemical elements are arranged according to their atomic

numbers, electron configurations and recurring chemical properties.

NAMES AND SYMBOLS OF THE FIRST TWENTY ELEMENTS

Recall that there are different kinds of elements. The table below shows the names and symbols

of the first twenty elements.

Name of

element

Chemical

symbol

Atomic number

(Z)

Number of

neutrons (N)

Atomic mass (Z + N)

Hydrogen H 1 0 1

Helium He 2 2 4

Lithium Li 3 4 7

Beryllium Be 4 5 9

Boron B 5 6 11

Carbon C 6 6 12

Nitrogen N 7 7 14

Oxygen O 8 8 16

Fluorine F 9 10 19

Neon Ne 10 10 20

Sodium Na 11 12 23

Magnesium Mg 12 12 24

Aluminium Al 13 14 27

Silicon Si 14 14 28

Phosphorous P 15 16 31

Sulphur S 16 16 32

Chlorine Cl 17 18 35

Argon Ar 18 22 40

Potassium K 19 20 39

Calcium Ca 20 20 40

Page | 43

When these elements are shown in the periodic table, their standard notations are used. For

example, hydrogen (H) is shown as

, magnesium (Mg) is shown as

, and so on.

The Periodic Table of the first twenty elements is shown below.

MAIN FEATURES OF THE PERIODIC TABLE

There are two main features of the periodic table: groups and periods.

GROUPS

These are vertical columns of elements in the periodic. Groups are also called families. There

eight groups in the periodic table. They are indicated by Roman numerals I, II, III, IV, V, VI,

VII, VII and VIII. For example,

 Group I contains hydrogen (H), lithium (Li), sodium (Na) and potassium (K).

 Group II contains beryllium (Be), magnesium (Mg) and calcium (Ca).

Elements in the same group:

 have the same number of electrons in the outermost shell.

 have similar physical and chemical properties.

 show trends in melting and boiling points

PERIODS

Page | 44

These are horizontal rows of elements in the periodic table. Periods are also called series. There

are four periods in the periodic table. They are indicated by Hindu Arabic numerals. For

example,

 Period 1 has hydrogen (H) and helium (He) only.

 Period 2 has lithium (Li), beryllium (Be), boron (B), carbon (C), nitrogen (N), oxygen (O),

fluorine (F) and neon (Ne).

All elements in the same period have the same number of electron shells.

GENERAL DISTRIBUTION OF ELEMENTS IN THE PERIODIC TABLE

The elements in the Periodic Table can be classified into metals, metalloids (semi-metals) and

non-metals.

 Metals (groups I, II and III).

 Metalloids (groups III and IV).

 Non-metals (groups V, VI, VII and VIII).

ELECTRON CONFIGURATION OF THE FIRST TWENTY ELEMENTS

Recall that electron configuration refers to the number and arrangement of electrons in an atom.

The Periodic Table below shows the electron configurations of the first twenty elements.

2.1

2.2

2.3

2.4

2.5

2.6

2.7

2.8

2.8.1

2.8.2

2.8.3

2.8.4

2.8.5

2.8.6

2.8.7

2.8.8

2.8.8.1

2.8.8.2

Page | 45

RELATING THE ELECTRON CONFIGURATION TO THE PERIODS AND GROUPS

OF THE PERIODIC TABLE

The electron configuration of any element can be used to predict the group number of an element

as well as the period to which the element belongs in the Periodic Table.

RELATIONSHIP BETWEEN VALENCE ELECTRONS AND GROUP NUMBER OF

THE ELEMENT

The number of electrons in the outermost shell (valence electrons) indicates the group number to

which the element belongs.

For example, all elements in group I have one valence electron, elements in group II have two

valence electrons, elements in group III have three valence electrons, and so on.

RELATIONSHIP BETWEEN PERIOD AND NUMBER OF SHELLS OF AN ATOM

The number of shells of an atom of an element indicates the period to which the element belongs.

For example all elements in period 1 have one electron shell only, all elements in period 2 have

two electron shells, elements in period 3 have three electron shells, and so on.

Example

An element X can be represented as:

a. Write down the electron configuration of X.

b. How many neutrons does the atom of X have?

c. To which group of the Periodic Table does element X belong? Give a reason.

d. To which period of the Periodic Table does element X belong? Give a reason.

SOLUTION

a. 2.8.8.1

b. Number of neutrons = Atomic mass – Atomic number

= 39 -19

= 20

Page | 46

c. It belongs to group 1 because it has one electron in its outermost shell.

d. It belongs to period 4 because it has four electron shells.

FAMILY NAMES OF ELEMENTS IN THE PERIODIC TABLE

Some of the groups in the Periodic Table have special names. These include:

1. Group I elements : Alkali metals

The alkali metals include Lithium (Li), Sodium (Na) and Potassium (K). Hydrogen is not an

alkali metal. It is a gaseous element. It is placed in group I mainly because it has one valence

electron just like all the alkali metals.

2. Group II elements : Alkaline earth metals

The alkaline earth metals include Beryllium (Be), Magnesium (Mg), and Calcium (Ca)

3. Group VII elements : Halogens

Halogens include Fluorine (F), Chlorine (Cl), Bromine (Br), and Iodine (I). Bromine and Iodine

are not shown in the Periodic Table of the first twenty elements but they are halogens.

4. Group VIII elements : Noble gases

Noble gases are also called inert gases. These include Helium (He), Neon (Ne) and Argon (Ar).

Page | 47

TOPIC 6 : PHYSICAL AND CHEMICAL CHANGES

A substance can be changed by heating it, mixing it with another substance, and adding water to

it, among other things. The change that takes place is either physical or chemical.

PHYSICAL CHANGE

A physical change is a process in which no new substance is formed. During a physical change,

only the physical properties of the substance are changed. Examples of physical changes are:

 Melting of candle wax

 Melting of ice into water

 Dissolving of sugar in water

 Changes of states of matter e.g. condensation, sublimation, freezing etc.

CHARACTERISTICS OF A PHYSICAL CHANGE

 No new substance is formed.

 No energy is either given out or absorbed in.

 The mass of the substance does not change.

 The change is usually reversible.

CHEMICAL CHANGE

A chemical change is a process in which a new substance is formed. Examples of chemical

changes are:

 combustion of fuel

 burning of wood, charcoal and paper.,

 chemical reactions

 decomposition of compounds

CHARACTERISTICS OF A CHEMICAL CHANGE

 A new substance is formed.

 Energy is usually given out or absorbed.

Page | 48

 The mass of substances changes.

 The change is usually irreversible.

DIFFERENCES BETWEEN PHYSICAL AND CHEMICAL CHANGES

Physical change Chemical change

No new substance is formed A new substance is formed

No energy is either given out or absorbed Energy is usually given out or absorbed

The mass of the substance does not change The mass of the substance changes

The change is usually reversible The change is usually irreversible

CHEMICAL REACTIONS

A chemical reaction is the re-arrangement of atoms to form new substances. In any chemical

reaction, there are two groups of substances: reactants and products.

Reactants

These are the substances that take part in and undergo change during a chemical reaction.

Products

These are substances that are formed as a result of the chemical reaction.

WHAT HAPPENS DURING A CHEMICAL REACTION?

 the bond in the reactant breaks

 the atoms in the reactant compounds form re-arrange or forms new bonds with other atoms.

This results in the formation of new substances.

CHEMICAL EQUATIONS

A chemical equation is a shorthand way of representing what happens in a chemical reaction. In

a chemical reaction:

 reactants and products are separated by an arrow as follows:

Reactants Products

Page | 49

For example:

Hydrogen + Oxygen Water

(Reactants) (Product)

NB: reactants are always written on the left hand side while products are written on the right

hand side of the arrow.

 the plus (+) sign in chemistry when used on the left hand side of the arrow means “reacts

with”.

 the arrow means “to form” the products on the right.

 a double arrow

( )

means that the reaction is reversible.

WAYS OF WRITING CHEMICAL EQUATIONS

Chemical equations can be written in two ways:

1. By using word equations

2. By using chemical symbols/formulae

WORD EQUATIONS

A word equation uses words to describe the reaction. In this equation you write the names of the

reactants and products. Examples are:

Magnesium + Oxygen Magnesium oxide

Sodium + Chlorine Sodium chloride

Calcium carbonate

heat

Calcium oxide + Carbon dioxide

CHEMICAL SYMBOLS/FORMULAE EQUATIONS

In this equation, you write the chemical formula of elements to show the reactants and products.

Example 1

Carbon reacts with oxygen to produce carbon dioxide. Write a chemical equation for this

reaction using chemical formulae.

Page | 50

SOLUTION

C + O

Example 2:

Magnesium reacts with chlorine to produce magnesium chloride. Write a chemical equation for

this reaction using chemical symbols.

SOLUTION

Mg + Cl

MgCl

BALANCING CHEMICAL EQUATIONS

Balancing a chemical equation is the process of ensuring that the same number of atoms of each

element is on either side of the equation.

GUIDELINES FOR BALANCING CHEMICAL EQUATIONS

1. Write the correct formulae for reactants and products.

2. Balance by multiplication of coefficients, not addition or subtraction of subscripts.

COEFFICIENT

This is a number that is written in front of a molecular formula. This number must be a whole

number, not a fraction.

SUBSCRIPT

This is a lowered small digit. It is written slightly below the symbol of the element or atom.

Page | 51

3. Write the physical states of the substance in the formula. The following are the short forms

for different physical states which are commonly used in chemical equations.

Physical state State symbol

solid state (s)

liquid state (l)

gaseous state (g)

aqueous solution (aq)

WORKED EXAMPLES ON BALANCING EQUATIONS

Example 1:

Balance the following equation: H

+ O

SOLUTION

Step 1:

Compare the number of atoms on the left hand side (L.H.S) and the right hand side (R.H.S).

Atom Left Hand Side Right Hand Side

Hydrogen (H) 2 2

Oxygen (O) 2 1

Step 2:

To make oxygen atoms equal, write numbers in front of the formula (H

). We can start with 2. If

does not balance, we go to 3, 4 until the equation is balanced.

+ O

When you have 2H

O, the number in front means that everything is multiplied by that number.

Thus, there are two hydrogen molecules and two oxygen atoms.

Step 3:

Page | 52

Count the number of atoms of each element on the reactants and products, we find that the total

number is the same.

Atom Left Hand Side Right Hand Side

Hydrogen (H) 4 4

Oxygen (O) 2 2

This means that the equation is now balanced.

Step 4:

Insert the correct state symbols for each substance.

(g) + O

(g) 2H

O (l)

CALCULATING MASSES OF REACTANTS AND PRODUCTS IN CHEMICAL

REACTIONS

To calculate the masses of reactants or products of a reaction follow the following simple steps:

 Write a balanced equation for the reaction

 Use the balanced equation to come up with the relative atomic masses (RAMs) of each

substance.

 Express the RAMs in grams.

 Use simple proportion to work out required masses.

Example 1:

Carbon reacts with oxygen to form carbon dioxide according to the equation below:

C (s) + O

(g) CO

(g)

(RAM of C = 12, O = 16)

Re-write the equation using masses.

SOLUTION

Page | 53

The relative atomic mass (RAM) of carbon, C = 12 amu

The relative atomic mass of oxygen, O = 16 amu

The molecular mass of oxygen molecule, O

= (16 + 16) amu = 32 amu

The molecular mass of carbon dioxide, CO

= (12 + 32) amu

= 44 amu

Expressing the masses in grams we have

12 g of carbon reacts with 32 g of oxygen to form 44 g of carbon dioxide. Using an equation:

12 g C + 32 g O 44 g CO

Example 2

Magnesium reacts with oxygen to form magnesium oxide according to the equation below:

2Mg (s) + O

(g) 2MgO (s)

If 12 g of magnesium is used:

a. how much oxygen is needed?

b. how much magnesium oxide is produced?

(RAM of Mg = 24, and O =16)

SOLUTION

From the equation:

Mass of magnesium = (2



24) g = 48 g

Mass of oxygen = (2



16) g = 32 g

Mass of magnesium oxide = 2



(24 +16) g = 80 g

Thus: 48 g Mg + 32 g O 80 g MgO

a. By proportion

Page | 54

48 g Mg = 32 g O

12 g Mg = y

y =

12 g × 32 g

48 g

= 8 g

Therefore, 8 g of oxygen is needed.

b. By proportion

48 g Mg = 80 g MgO

12 g Mg = x

x =

12 g× 80 g

48 g

= 20 g

Therefore, 20 g of magnesium oxide is produced.

THE LAW OF CONSERVATION OF MATTER

The law states that matter is neither created nor destroyed in a chemical reaction.

This statement means that in any chemical reaction, the total mass of the reactants is always

equal to the total mass of the products. For example when hydrogen reacts with oxygen to

produce water:

(g) + O

(g) 2H

O (l)

4 g H + 32 g O = 36 g H

36 g (reactants) = 36 g (product)

PERCENTAGE COMPOSITION OF A COMPOUND

The percentage composition of a compound tells us which elements are in the compound and

how much of each there is, expressed as a percentage of the total mass.

Page | 55

CALCULATING PERCENTAGE COMPOSITION BY MASS OF ELEMENTS IN A

COMPOUND

To calculate the percentage composition by mass of an element in a compound, follow the

following steps:

1. Write down the formula of the compound

2. Work out its relative formula mass (R.F.M) using the relative atomic masses (RAM) of each

element

3. Write the mass of the element in question as a fraction of the total.

4. Multiply the fraction by 100.

Example 1:

Calculate the percentage composition of each element in calcium carbonate (CaCO

). (RAM of

Ca = 40, C =12 and O = 16).

SOLUTION

Relative Formula Mass of CaCO

= (40



1) + (12



1) + (16



= (40 + 12 + 48) amu

= 100 amu

a. Percentage composition of calcium (Ca) = mass of calcium



100

RFM of CaCO

40 amu

×100

100 amu

= 40%

b. Percentage composition of oxygen (O) = mass of oxygen



100

RFM of CaCO

(16×3) amu

×100

100amu

48 amu

×100

100 amu

Page | 56

= 48%

c. Percentage composition of carbon (C) = mass of carbon



100

RFM of CaCO

12 amu

×100

100 amu

= 12%

Example 2:

Work out the percentage composition by mass of each element in ammonium nitrate (NH

(RAM of N = 14, H = 1 and O =16).

SOLUTION

Relative formula mass of NH

(14 2 ) (1 4) (16 3)    

amu

= (28 + 4 + 48) amu

= 80 amu

a. Percentage composition of nitrogen (N) = mass of nitrogen



100

RFM of NH

(14 × 2) amu

×100

80 amu

28 amu

×100

80 amu

= 35%

b. Percentage composition of hydrogen (H) = mass of hydrogen



100

RFM of NH

(1× 4) amu

×100

80 amu

4 amu

×100

80 amu

Page | 57

= 5%

c. Percentage composition of oxygen (O) = mass of oxygen



100

RFM of NH

(16 × 3) amu

×100

80 amu

48amu

×100

80 amu

= 60%

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TOPIC 7 : ORGANIC COMPOUNDS

An organic compound is a compound that contains the element carbon in its molecule. A branch

of chemistry that deals with organic compounds is called organic chemistry. Examples of

organic compounds are proteins, carbohydrates, fats, crude oil, plastics and medical drugs made

from crude oil.

Compounds such as carbon monoxide, carbon dioxide, metal carbonates and hydrogen

carbonates are not considered to be organic compounds even though they contain carbon in their

molecules.

HISTORY OF ORGANIC COMPOUNDS

A long time ago, scientists believed that organic compounds could only be made by natural

living things. For example compounds such as Deoxyribonucleic acid (DNA) and insulin found

in human bodies are examples of organic compounds made by living things. However, now it is

a known fact that this is not true. Apart from organic compounds deriving from living things,

they can also be made artificially by organic chemists.

SOURCES OF ORGANIC COMPOUNDS

1. Plants and animals

These organisms synthesize many organic compounds. The compounds include sugar, starch,

fats, dyes, drugs among other things.

2. Fossil fuels

These are fuels derived from the decomposed remains of plants and animals which once lived

and died in the sea.

3. Natural gas

Natural gas is a hydrocarbon found underground on top of crude oil and under the sea. It is

95% methane.

4. Coal

This is a fossil fuel formed from decayed plant materials that has been subjected to heat and

pressure over millions of years.

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ORGANIC COMPOUNDS AS FUELS

A fuel is a substance that is used to provide energy. A fuel can be in form of solid, liquid or gas.

SUBSTANCES USED AS FUELS IN HOMES

 Charcoal

 Butane

 Paraffin

 Wood

 Coal

 Petrol

 Diesel

 Ethanol

 Methylated spirit (ethanol mixed with methanol in spirit lamps and stoves)

CLASSES OF FUELS

There are two classes of fuels: bio fuels and fossil fuels.

1. BIO FUELS

These are fuels made from plants that grow around us. Bio fuels are renewable i.e. they can be

replaced after we have used them by growing more plants. Examples of bio fuels are wood,

biogas, biodiesel and ethanol. These fuels are made from plant materials. For example, ethanol

(commonly called alcohol) and biodiesel can be made from starch and sugar.

2. FOSSIL FUELS

These are fuels formed from remains of decayed plants and animals. They are found in the

ground. Examples of fossil fuels are: petroleum or crude oil, coal and natural gas.

PETROLEUM

Petroleum is a mixture of hydrocarbons. A hydrocarbon is a compound that contains carbon and

hydrogen atoms in its molecule. Petroleum is also called crude oil.

Page | 60

COMPOSITION OF PETROLEUM

The components of petroleum are called fractions. These include:

 Petrol

 Diesel

 Bitumen

 Paraffin

 Naphtha

 Lubricants

SEPARATION OF COMPONENTS OF PETROLEUM

The components are separated by fractional distillation. This is because the fractions have

different boiling points. The crude oil is evaporated and its vapours allowed to condense at

different temperatures in the fractionating column.

Page | 61

USES OF FRACTIONS OF PETROLEUM

The following are some of the uses of the fractions of petroleum.

Fraction Uses

Petro (gasoline)



It is used as a fuel for vehicles

Kerosene (paraffin)



Used as jet fuel

 Use for heating and lighting

Diesel



Used as fuel for heavy vehicles

Bitumen



Used for tarmacking roads

 Used for repairing leaking roofs

Lubricants (e.g. lubricating oil or grease)



Used for lubricating moving parts of

machines

 Used for making waxes

 Used for making polishes